Coordination Compounds: NEET notes

1. Werner's theory

NCERT §5.1

• A coordination compound holds a central metal atom or ion joined to a set of anions or neutral molecules through shared electron pairs. Chlorophyll (Mg), haemoglobin (Fe) and vitamin B12 (Co) are all coordination compounds.
• Werner tested the cobalt(III) chloride-ammonia compounds with excess cold AgNO₃. Per mole: CoCl₃·6NH₃ (yellow) gave 3 mol AgCl, CoCl₃·5NH₃ (purple) gave 2, and both forms of CoCl₃·4NH₃ (green and violet) gave 1.
• He explained this by keeping six groups (Cl⁻, NH₃ or both) fixed on the cobalt in a unit that does not break up: [Co(NH₃)₆]Cl₃ is a 1:3 electrolyte, [CoCl(NH₃)₅]Cl₂ a 1:2 electrolyte and [CoCl₂(NH₃)₄]Cl a 1:1 electrolyte, as their conductivities confirm.
• Only chloride outside the square bracket is free to meet Ag⁺; chloride bonded to the metal inside the bracket is not precipitated.
• The green and violet CoCl₃·4NH₃ share one empirical formula but differ in properties, so they are isomers.
• Werner's postulates (1898): a metal shows primary and secondary valences. Primary valences are ionisable and satisfied by negative ions. Secondary valences are non-ionisable, satisfied by neutral molecules or anions, equal to the coordination number and fixed for a given metal.
• The groups held by secondary valences have a definite arrangement in space for each coordination number, today called the coordination polyhedron. The bracketed species is the complex; ions outside are counter ions.
• Common shapes: [Co(NH₃)₆]³⁺, [CoCl(NH₃)₅]²⁺ and [CoCl₂(NH₃)₄]⁺ are octahedral, [Ni(CO)₄] is tetrahedral and [PtCl₄]²⁻ is square planar.
• A double salt such as carnallite (KCl·MgCl₂·6H₂O), Mohr's salt (FeSO₄·(NH₄)₂SO₄·6H₂O) or potash alum breaks up completely into simple ions in water. A complex does not: K₄[Fe(CN)₆] gives [Fe(CN)₆]⁴⁻, not Fe²⁺ and CN⁻.

2. Ligands and denticity

NCERT §5.2

• A coordination entity is one metal atom or ion with a set number of ions or molecules attached to it. Examples: [CoCl₃(NH₃)₃], [Ni(CO)₄], [PtCl₂(NH₃)₂], [Fe(CN)₆]⁴⁻ and [Co(NH₃)₆]³⁺.
• The central atom or ion holds the ligands in a definite geometry and accepts their electron pairs, so it acts as a Lewis acid: Ni²⁺ in [NiCl₂(H₂O)₄], Co³⁺ in [CoCl(NH₃)₅]²⁺ and Fe³⁺ in [Fe(CN)₆]³⁻.
• Ligands are the ions or molecules bonded to the central atom. They range from simple ions (Cl⁻) and small molecules (H₂O, NH₃) to larger molecules and even proteins.
• Unidentate ligands bind through one donor atom (Cl⁻, H₂O, NH₃). Didentate ligands bind through two: ethane-1,2-diamine (en, H₂NCH₂CH₂NH₂) and oxalate (C₂O₄²⁻). Polydentate ligands, such as N(CH₂CH₂NH₂)₃, have several.
• EDTA⁴⁻ (ethylenediaminetetraacetate) is hexadentate: it binds one metal ion through two nitrogen and four oxygen atoms.
• A chelate ligand uses two or more donor atoms on the same metal ion at once, and the number of such donor groups is its denticity. Chelate complexes are generally more stable than similar complexes of unidentate ligands.
• An ambidentate ligand has two different donor atoms and binds through either one: NO₂⁻ through N or O, and SCN⁻ through S or N.

3. Coordination number and oxidation number

NCERT §5.2

• The coordination number (CN) is the number of ligand donor atoms bonded directly to the metal: 6 for Pt in [PtCl₆]²⁻ and 4 for Ni in [Ni(NH₃)₄]²⁺.
• A didentate ligand counts twice: Fe in [Fe(C₂O₄)₃]³⁻ and Co in [Co(en)₃]³⁺ both have CN = 6.
• Only σ bonds from ligand to metal fix the coordination number; any π bonds between them are not counted.
• The coordination sphere is the central atom with its ligands, written in square brackets. Ionisable groups outside the bracket are counter ions: in K₄[Fe(CN)₆] the sphere is [Fe(CN)₆]⁴⁻ and the counter ion is K⁺.
• The coordination polyhedron is the spatial arrangement of the donor atoms around the central atom. The usual ones are octahedral ([Co(NH₃)₆]³⁺), square planar ([PtCl₄]²⁻) and tetrahedral ([Ni(CO)₄]).
• The oxidation number of the central atom is the charge left on it if every ligand is taken away together with the electron pair it shares. It is written as a Roman numeral in brackets: Cu in [Cu(CN)₄]³⁻ is +1, Cu(I).
• Homoleptic complexes carry one kind of donor group, e.g. [Co(NH₃)₆]³⁺; heteroleptic complexes carry more than one, e.g. [Co(NH₃)₄Cl₂]⁺.

4. Formulas and names

NCERT §5.3

• Formula rules: write the central atom first, then the ligands in alphabetical order whatever their charge. Put the whole entity in square brackets, polyatomic ligands and abbreviations in round brackets, and leave no spaces inside the sphere.
• An abbreviated ligand is alphabetised by the first letter of its abbreviation. For a charged entity written alone, the charge goes outside the bracket as a right superscript with the number before the sign: [Co(CN)₆]³⁻, [Cr(H₂O)₆]³⁺.
• Naming rules: name the cation first, whether or not it is the complex. Inside the entity, name the ligands in alphabetical order before the metal, which is the reverse of the order used when writing the formula.
• Anionic ligands end in -o (the 2004 IUPAC draft makes this -ido, so chloro becomes chlorido). Neutral ligands keep their names except aqua (H₂O), ammine (NH₃), carbonyl (CO) and nitrosyl (NO).
• Use di, tri, tetra for numbers of simple ligands. When a ligand's own name already holds a number, use bis, tris, tetrakis with the ligand in brackets: [NiCl₂(PPh₃)₂] is dichloridobis(triphenylphosphine)nickel(II).
• Give the metal's oxidation state as a Roman numeral in brackets. In a cationic or neutral complex the metal keeps its element name; in an anionic complex it ends in -ate (cobaltate), with Latin roots for some metals (ferrate for Fe, argentate for Ag).
• Worked names: [Cr(NH₃)₃(H₂O)₃]Cl₃ is triamminetriaquachromium(III) chloride; [Co(en)₃]₂(SO₄)₃ is tris(ethane-1,2-diamine)cobalt(III) sulphate; [Ag(NH₃)₂][Ag(CN)₂] is diamminesilver(I) dicyanidoargentate(I).
• Finding the oxidation state: charge on the entity = oxidation state + sum of ligand charges. In [Cr(NH₃)₃(H₂O)₃]³⁺ all ligands are neutral, so Cr is +3; in K₃[Fe(CN)₆], −3 = x + 6(−1) gives Fe = +3.
• Name endings show binding atoms for ambidentate ligands: nitrito-N (through N) and nitrito-O (through O); thiocyanato-S binds through S, as in Hg[Co(SCN)₄], mercury(I) tetrathiocyanato-S-cobaltate(III).

5. Geometrical and optical isomerism

NCERT §5.4.1; §5.4.2

• Isomers share a formula but differ in how the atoms are arranged. Stereoisomers have the same bonds in different spatial arrangements; structural isomers have different bonds.
• Geometrical isomerism appears in heteroleptic complexes of CN 4 and 6. In square planar [MX₂L₂], the two X ligands sit side by side (cis) or across from each other (trans), as in cis- and trans-[Pt(NH₃)₂Cl₂].
• Square planar [MABXL] with four different unidentate ligands has three geometrical isomers: two cis and one trans.
• Tetrahedral complexes never show geometrical isomerism: every position is adjacent to every other, so the ligands' relative positions are all alike.
• Octahedral [MX₂L₄] ([Co(NH₃)₄Cl₂]⁺) and [MX₂(L–L)₂] ([CoCl₂(en)₂]⁺) also form cis and trans isomers.
• Octahedral [Ma₃b₃], such as [Co(NH₃)₃(NO₂)₃], gives fac and mer isomers: three identical donors on the corners of one face (facial), or along a meridian of the octahedron (meridional).
• Optical isomers (enantiomers) are non-superimposable mirror images, and such molecules or ions are chiral. The dextro (d) form rotates plane-polarised light to the right and the laevo (l) form to the left in a polarimeter.
• Optical isomerism is common in octahedral complexes with didentate ligands, e.g. [Co(en)₃]³⁺. For [PtCl₂(en)₂]²⁺ only the cis isomer is optically active; the trans form has a mirror plane.
• Of cis- and trans-[CrCl₂(ox)₂]³⁻, only the cis form is chiral.

6. Structural isomerism

NCERT §5.4.3–5.4.6

• Linkage isomerism needs an ambidentate ligand. Thiocyanate may bind as M–NCS or M–SCN.
• Jørgensen found [Co(NH₃)₅(NO₂)]Cl₂ in two forms: red, with nitrite bound through oxygen (–ONO), and yellow, with nitrite bound through nitrogen (–NO₂).
• Coordination isomerism arises when ligands swap between the cationic and anionic entities of two different metals: [Co(NH₃)₆][Cr(CN)₆] and [Cr(NH₃)₆][Co(CN)₆].
• Ionisation isomerism arises when a counter ion that could itself be a ligand trades places with a ligand: [Co(NH₃)₅(SO₄)]Br and [Co(NH₃)₅Br]SO₄.
• Ionisation isomers give different ions in water, so they answer different tests: [Co(NH₃)₅Br]SO₄ gives white BaSO₄ with Ba²⁺, while [Co(NH₃)₅(SO₄)]Br gives AgBr with Ag⁺.
• Solvate isomerism (hydrate isomerism when the solvent is water) depends on whether solvent is bonded to the metal or just sits in the crystal: [Cr(H₂O)₆]Cl₃ is violet, while [Cr(H₂O)₅Cl]Cl₂·H₂O is grey-green.

7. Valence bond theory

NCERT §5.5.1

• Werner's theory could not explain why only some elements form complexes, why the bonds are directional, or where the magnetic and optical properties come from. VBT and CFT are the two models treated at this level.
• In VBT the metal mixes (n−1)d, ns, np orbitals, or ns, np, nd orbitals, to make a set of identical hybrid orbitals pointing in fixed directions. Each hybrid orbital accepts an electron pair from a ligand.
• Hybridisation and shape: CN 4, sp³ tetrahedral; CN 4, dsp² square planar; CN 5, sp³d trigonal bipyramidal; CN 6, sp³d² or d²sp³ octahedral.
• [Co(NH₃)₆]³⁺: Co³⁺ is 3d⁶. The six electrons pair into three 3d orbitals, leaving two inner 3d orbitals for d²sp³ hybridisation. No electron is unpaired, so the octahedral complex is diamagnetic: an inner orbital, low spin or spin paired complex.
• [CoF₆]³⁻ uses the outer 4d orbitals (sp³d²), keeps four unpaired electrons and is paramagnetic: an outer orbital, high spin or spin free complex.
• [NiCl₄]²⁻: Ni²⁺ is 3d⁸, the four Cl⁻ pairs go into sp³ hybrids, the complex is tetrahedral and two unpaired electrons make it paramagnetic. [Ni(CO)₄] is also tetrahedral but diamagnetic, because nickel is at zero oxidation state with no unpaired electron.
• [Ni(CN)₄]²⁻: the 3d⁸ electrons pair up to free one 3d orbital, giving dsp² hybrids. The complex is square planar and diamagnetic.
• Hybrid orbitals are not real objects; hybridisation is a mathematical treatment of the atomic wave functions.
• Magnetic data often reveal the geometry in VBT: a paramagnetic four-coordinate Ni²⁺ complex is tetrahedral, while a diamagnetic one is square planar.

8. Magnetic properties and limits of VBT

NCERT §5.5.2; §5.5.3

• Magnetic susceptibility measurements give the magnetic moment, hence the number of unpaired electrons and a clue to the structure.
• d¹ to d³ ions (Ti³⁺, V³⁺, Cr³⁺) always have two vacant 3d orbitals for octahedral hybridisation, so the free ion and its complexes behave alike magnetically.
• For d⁴ (Cr²⁺, Mn³⁺), d⁵ (Mn²⁺, Fe³⁺) and d⁶ (Fe²⁺, Co³⁺), two 3d orbitals can be emptied only by pairing, which leaves 2, 1 and 0 unpaired electrons respectively.
• Magnetic data show both behaviours: [Mn(CN)₆]³⁻ has 2 unpaired electrons but [MnCl₆]³⁻ has 4; [Fe(CN)₆]³⁻ has 1 but [FeF₆]³⁻ has 5; [Co(C₂O₄)₃]³⁻ is diamagnetic but [CoF₆]³⁻ has 4.
• VBT's answer: [Mn(CN)₆]³⁻, [Fe(CN)₆]³⁻ and [Co(C₂O₄)₃]³⁻ are inner orbital (d²sp³) complexes; [MnCl₆]³⁻, [FeF₆]³⁻ and [CoF₆]³⁻ are outer orbital (sp³d²) complexes.
• Example: [MnBr₄]²⁻ has a spin-only moment of 5.9 BM, meaning five unpaired electrons, so Mn²⁺ has not paired any d electrons: the ion is tetrahedral (sp³), not square planar (dsp²).
• Limits of VBT: it rests on several assumptions, gives no quantitative account of magnetic data or of thermodynamic and kinetic stability, says nothing about colour, cannot reliably choose between tetrahedral and square planar for CN 4, and does not separate weak from strong ligands.

9. Crystal field theory

NCERT §5.5.4

• CFT treats the metal-ligand bond as purely ionic: anionic ligands are point charges and neutral ligands are point dipoles, and they interact with the metal only electrostatically.
• In a free metal ion, or in a spherical field of charge, the five d orbitals have equal energy (degenerate). An asymmetric field from real ligands lifts this degeneracy: crystal field splitting.
• Octahedral field: dx²−y² and dz² point at the ligands, feel more repulsion and rise as the eg pair; dxy, dyz and dxz point between the axes and fall as the t₂g set. The gap is Δo.
• Relative to the average energy, the two eg orbitals go up by (3/5)Δo and the three t₂g orbitals go down by (2/5)Δo.
• Δo grows with the field of the ligand and with the charge on the metal ion. The spectrochemical series, found by experiment from absorption spectra, orders ligands by field strength: I⁻ < Br⁻ < SCN⁻ < Cl⁻ < S²⁻ < F⁻ < OH⁻ < C₂O₄²⁻ < H₂O < NCS⁻ < edta⁴⁻ < NH₃ < en < CN⁻ < CO.
• d¹ to d³ ions put their electrons singly into t₂g (Hund's rule). For d⁴ the fourth electron has a choice, set by Δo against the pairing energy P.
• If Δo < P the fourth electron goes into eg, giving t₂g³ eg¹: weak field ligands, high spin complexes. If Δo > P it pairs in t₂g, giving t₂g⁴ eg⁰: strong field ligands, low spin complexes.
• Calculations show that complexes of d⁴ to d⁷ ions are more stable in a strong field than in a weak one.
• Tetrahedral field: the splitting is inverted and smaller, Δt = (4/9)Δo for the same metal, ligands and distances. It is rarely large enough to force pairing, so tetrahedral complexes are almost always high spin.
• The g subscript (t₂g, eg) is used only for complexes with a centre of symmetry, such as octahedral and square planar ones; tetrahedral levels are written without it.

10. Colour in coordination compounds

NCERT §5.5.5; §5.5.6

• A coloured complex absorbs part of the visible light passing through it; its colour is the complementary colour of what is left. If green is absorbed, the complex looks red.
• Absorbed wavelength and colour seen: [CoCl(NH₃)₅]²⁺ 535 nm (yellow) looks violet; [Co(NH₃)₅(H₂O)]³⁺ 500 nm (blue-green) looks red; [Co(NH₃)₆]³⁺ 475 nm (blue) looks yellow-orange.
• More pairs: [Co(CN)₆]³⁻ absorbs at 310 nm in the ultraviolet and is pale yellow; [Cu(H₂O)₄]²⁺ absorbs 600 nm (red) and looks blue; [Ti(H₂O)₆]³⁺ absorbs 498 nm (blue-green) and looks violet.
• [Ti(H₂O)₆]³⁺ is d¹: the lone electron sits in t₂g. Blue-green light promotes it to eg (t₂g¹eg⁰ → t₂g⁰eg¹), so the complex looks violet. CFT puts colour down to such d-d transitions.
• No ligand, no splitting, no colour: heating [Ti(H₂O)₆]Cl₃ to drive off water leaves a colourless solid, and anhydrous CuSO₄ is white while CuSO₄·5H₂O is blue.
• Changing the ligand changes the colour. Adding en to green [Ni(H₂O)₆]²⁺ in en:Ni ratios 1:1, 2:1 and 3:1 gives pale blue [Ni(H₂O)₄(en)]²⁺, blue/purple [Ni(H₂O)₂(en)₂]²⁺ and violet [Ni(en)₃]²⁺.
• Ruby is Al₂O₃ with about 0.5-1% Cr³⁺ (d³) in some Al³⁺ sites; these act like octahedral Cr(III) complexes whose d-d transitions give the colour. In emerald, Cr³⁺ sits in octahedral sites of beryl (Be₃Al₂Si₆O₁₈); its bands move to longer wavelengths (yellow-red and blue), so green is transmitted.
• Limits of CFT: treating ligands as point charges predicts that anions should split most, yet anionic ligands sit at the weak end of the spectrochemical series; and CFT ignores the covalent side of metal-ligand bonding. Ligand field and molecular orbital theories handle these.

11. Metal carbonyls

NCERT §5.6

• Homoleptic carbonyls contain only CO ligands; most transition metals form them, with simple well-defined structures.
• Shapes: tetracarbonylnickel(0), Ni(CO)₄, is tetrahedral; pentacarbonyliron(0), Fe(CO)₅, is trigonal bipyramidal; hexacarbonylchromium(0), Cr(CO)₆, is octahedral.
• Decacarbonyldimanganese(0), [Mn₂(CO)₁₀], is built from two Mn(CO)₅ square pyramids held together by one metal-metal bond. In octacarbonyldicobalt(0), [Co₂(CO)₈], the two cobalt atoms are bonded directly and two CO groups also bridge them.
• The M–C bond has both σ and π character. The σ bond forms when the carbon's lone pair is donated into an empty metal orbital.
• The π bond forms when a filled metal d orbital donates electrons back into the empty antibonding π* orbital of CO (back bonding).
• The two donations reinforce each other (synergic bonding): σ donation makes the metal richer in electrons for back donation, and back donation lets it accept more σ density. This strengthens the metal-CO bond.

12. Importance and applications

NCERT §5.7

• Analysis: colour reactions of metal ions with chelating ligands form the basis of their detection and estimation. Reagents include EDTA, DMG (dimethylglyoxime), α-nitroso-β-naphthol and cupron.
• Hardness of water is found by titration with Na₂EDTA, which forms stable complexes with Ca²⁺ and Mg²⁺; their different stability constants let the two ions be estimated selectively.
• Metallurgy: with cyanide, air (O₂) and water, gold dissolves as the complex [Au(CN)₂]⁻, and adding zinc later recovers the metal. Silver is extracted in a similar way.
• Purification: impure nickel is turned into [Ni(CO)₄], which is then decomposed to give pure nickel.
• Biology: chlorophyll is a magnesium complex, haemoglobin (the oxygen carrier in blood) an iron complex, and vitamin B12 (cyanocobalamine, the anti-pernicious anaemia factor) a cobalt complex. The enzymes carboxypeptidase A and carbonic anhydrase also hold coordinated metal ions.
• Industry: Wilkinson's catalyst, the rhodium complex [(Ph₃P)₃RhCl], hydrogenates alkenes.
• Electroplating from [Ag(CN)₂]⁻ and [Au(CN)₂]⁻ gives smoother, more even silver and gold coats than simple metal ions do.
• Black and white photography: hypo fixes the developed film by dissolving the unexposed AgBr as the complex [Ag(S₂O₃)₂]³⁻.
• Chelate therapy: D-penicillamine and desferrioxime B remove excess copper and iron, EDTA treats lead poisoning, and cis-platin and related platinum complexes inhibit the growth of tumours.

Must-know facts

1. Werner: 1 mol each of CoCl₃·6NH₃, ·5NH₃ and ·4NH₃ give 3, 2 and 1 mol AgCl; they are 1:3, 1:2 and 1:1 electrolytes.
2. Primary valence = ionisable = oxidation state; secondary valence = non-ionisable = coordination number.
3. Double salts (Mohr's salt, carnallite, potash alum) give all their simple ions in water; complexes such as K₄[Fe(CN)₆] do not.
4. EDTA⁴⁻ is hexadentate (2 N + 4 O); en and oxalate are didentate; NO₂⁻ and SCN⁻ are ambidentate.
5. Coordination number counts σ-bonded donor atoms: [Co(en)₃]³⁺ and [Fe(C₂O₄)₃]³⁻ both have CN 6.
6. Formula: metal first, ligands alphabetical. Name: ligands alphabetical, then metal; anionic complex ends in -ate (ferrate, argentate, cobaltate).
7. aqua = H₂O, ammine = NH₃, carbonyl = CO, nitrosyl = NO; anionic ligands end in -ido in the 2004 IUPAC draft (chlorido, cyanido).
8. Tetrahedral complexes show no geometrical isomerism; square planar [MABXL] has three geometrical isomers.
9. [Co(NH₃)₃(NO₂)₃] has fac and mer isomers; [Co(en)₃]³⁺ is optically active; only cis-[PtCl₂(en)₂]²⁺ is chiral.
10. Linkage: [Co(NH₃)₅(ONO)]²⁺ red, [Co(NH₃)₅(NO₂)]²⁺ yellow. Hydrate: [Cr(H₂O)₆]Cl₃ violet, [Cr(H₂O)₅Cl]Cl₂·H₂O grey-green.
11. [Co(NH₃)₆]³⁺: d²sp³, inner orbital, diamagnetic. [CoF₆]³⁻: sp³d², outer orbital, 4 unpaired.
12. [NiCl₄]²⁻: sp³, tetrahedral, 2 unpaired. [Ni(CN)₄]²⁻: dsp², square planar, diamagnetic. [Ni(CO)₄]: tetrahedral, diamagnetic, Ni(0).
13. [Fe(CN)₆]³⁻ has 1 unpaired electron, [FeF₆]³⁻ has 5; [Mn(CN)₆]³⁻ has 2, [MnCl₆]³⁻ has 4.
14. Octahedral splitting: eg +(3/5)Δo, t₂g −(2/5)Δo; tetrahedral Δt = (4/9)Δo, inverted, nearly always high spin.
15. Spectrochemical series ends: I⁻ weakest ... en < CN⁻ < CO strongest; H₂O < NH₃ < en.
16. Δo < P: high spin t₂g³eg¹ (weak field); Δo > P: low spin t₂g⁴eg⁰ (strong field).
17. [Ti(H₂O)₆]³⁺ absorbs 498 nm (blue-green) and looks violet; [Cu(H₂O)₄]²⁺ absorbs 600 nm (red) and looks blue.
18. Ruby: Al₂O₃ with 0.5-1% Cr³⁺; emerald: Cr³⁺ in beryl, Be₃Al₂Si₆O₁₈.
19. Carbonyls: Ni(CO)₄ tetrahedral, Fe(CO)₅ trigonal bipyramidal, Cr(CO)₆ octahedral; bonding is synergic σ donation + π back donation.
20. Wilkinson's catalyst [(Ph₃P)₃RhCl] hydrogenates alkenes; cis-platin fights tumours; EDTA treats lead poisoning; hypo gives [Ag(S₂O₃)₂]³⁻.

Common traps

Formulas

Key terms

Coordination entity

A central metal atom or ion together with the fixed set of ligands bonded to it.

Ligand

An ion or molecule that gives an electron pair to the central atom in a complex.

Denticity

The number of donor atoms through which one ligand binds a single metal ion.

Chelate

A complex in which one ligand grips the same metal through two or more donor atoms, forming a ring.

Ambidentate ligand

A ligand with two different donor atoms that can bind through either, such as NO₂⁻ or SCN⁻.

Coordination number

The number of ligand donor atoms σ-bonded directly to the central atom.

Coordination sphere

The central atom and its ligands, written inside square brackets.

Counter ion

An ionisable ion written outside the square bracket that balances the complex's charge.

Coordination polyhedron

The shape traced by the donor atoms around the central atom, such as an octahedron.

Homoleptic / heteroleptic

Having only one kind of ligand / having more than one kind.

Primary valence

Werner's ionisable valence, satisfied by anions; today the oxidation state.

Secondary valence

Werner's non-ionisable valence; today the coordination number.

Inner / outer orbital complex

A complex that hybridises (n−1)d orbitals (d²sp³, low spin) / nd orbitals (sp³d², high spin).

Crystal field splitting (Δo)

The energy gap between the eg and t₂g orbitals created by six ligands in an octahedron.

Pairing energy (P)

The energy cost of putting a second electron into an orbital that already holds one.

Spectrochemical series

Ligands ranked by the size of the d-orbital splitting they cause, found from absorption spectra.

Synergic bonding

σ donation from CO to the metal and π back donation from metal to CO, each strengthening the other.